Chapter 19 - Equilibrium

homogenous equilibrium

the equilibrium species all have the same state or phase

heterogenous equilibrium

the equilibrium species have different states or phases

-concentration of solids and liquids do not vary so is only gases or aqueous states included

how to find Kc from moles

R - molar ratio

I - initial moles

C - change in moles

E - equilibrium moles

find conc with equilibrium moles then find Kc

Kp

equilibrium constant for gases in terms of partial pressures

mole fraction

the proportion by volume of a gas to the total volume of gases in a gas mixture

-the sum of the mole fractions in a gas must equal 1

mole fraction EQUATION

x = number of moles of gas / total number of moles in gas mixture

partial pressure

the contribution that the gas makes towards the total pressure

partial pressure EQUATION

mole fraction x total pressure

p = x x P

Kp expression

p(products)^{no. of moles} / p(reactants)^{no. of moles}

round brackets = partial pressures

units = Pa, kPa or atm

value of K

-if K=1 indicates an equilibrium halfway between reactants and products

-if K near to 100 indicates equilibrium lies close to products - most reactants have turned into products

-if K near to 0.01 indicates equilibrium lies close to reactants - most reactants have not reacted

the effect of increasing TEMPERATURE -exothermic

-the position of equilibrium shifts to the left

-so there is a smaller concentration/partial pressure for products and greater concentration/partial pressure for reactants

-so value of K will decrease

the effect of increasing TEMPERATURE -endothermic

-the position of equilibrium shifts to the right

-so there is a greater concentration/partial pressure for products and smaller concentration/partial pressure for reactants

-so value of K will increase

the effect of CONC + PRESSURE

does not change equilibrium constant as Le Chatelier’s Principle causes position of equilibrium to shift to keep the constant the same

CONC example: A +2B ⇌ C + D - conc of C decreased

-Kc must remain constant

-C must be increased to counteract so D is also increased

-so A + B are decreased as more reactants turned into products

-so equilibrium position will shift to right

PRESSURE example: N₂O₄ ⇌ 2NO₂ -total pressure doubled

-Kp must remain constant

-partial pressure of NO2 decreases and partial pressure of N2O4 increases

-so equilibrium position will shift to the left - side with fewer moles

the effect of a CATALYST

-increases both forward + reverse reactions by the same amount

-the rate at which equilibrium constant is reached is quicker but position of equilibrium is unchanged